d – AND f – BLOCK ELEMENTS

TINTO JOHNS M. Sc., B. Ed

MELTING POINT AND BOILING POINT

• High M.P and B.P - Due to strong metallic bond and the presence of half filled d- orbitals

• Involvement of greater number of electrons from (n-1)d in addition to the ns electrons in the inter atomic metallic bonding.

• Because of stronger interatomic bonding, transition elements have high M.P and B.P

• In moving along the period from left to right, the M.P of these metals first INCREASES to MAXIMUM and the DECREASES regularly towards the end of the period.

• melting points of these metals rise to a maximum at

d5 except for anomalous values of Mn and Tc and fall regularly as the atomic number increases.TRENDS OF M.P OF 3- d , 4-d AND 5-d TRANSITION

METALS• The strength of interatomic bonds in transition

elements is roughly related to the number of half filled d- orbitals

• In the beginning the no. of half filled d- orbitals increases till the middle of the period causing increase in strength of interparticle bonds But thereafter the pairing of electrons in d – orbitals occurs and the no. of half filled orbitals decreases , which also cause deacrease in M.P

Trends in enthalpies of atomization of transition elements

1. greater the number of valence electrons, stronger the inter atomic attraction, hence stronger bonding between atoms resulting in higher enthalpies of atomization.

2. metals of the second and third series have greater enthalpies of atomization than the corresponding elements of the first series

Atomic and ionic radii• The Atomic/ionic radii first DECREASES till the

middle, becomes almost constant and then INCREASES towards the end of the period.

• New electron enters a d orbital each time the nuclear charge increases by unity, But the shielding effect of a d electron is not that effective, hence the net electrostatic attraction between the nuclear charge and the outermost electron increases and the ionic radius decreases

• However the increased nuclear charge is partly cancelled by the increased screening effect of electrons in the d – orbitals of penultimate shell.

• When the increased nuclear charge and increased Screening effect balance each other, the atomic radii becomes almost constant.

• Increase in atomic radii towards the end may be attributed to the electron – electron repulsion.

• In fact the pairing of electrons in d – orbitals occurs after d5 configuration.

• The repulsive interaction between the paired electron causes Increase in Atomic/ ionic radii

Sc = [Ar]4s2 3d1

Ti = [Ar]4s2 3d2

3d< 4d= 5d

• There is increase from the first (3d) to the second (4d) series of the elements.

• But the radii of the third (5d) series are virtually the same as 4d

• This is due to the intervention of the 4f orbital which must be filled before the 5d series of elements begin.

• There is a steady decrease in atomic radii from La due to the poor shielding of inner core electrons (4f) is known lanthanoid contraction.

IONISATION ENTHALPIES• Due to an increase in nuclear charge there is an

increase in ionisation enthalpy along each series of the transition elements from left to right.

• Ionisation enthalpies give some guidance concerning the relative stabilities of oxidation states.

• Although the first ionisation enthalpy, in general, increases, the magnitude of the increase in the second and third ionisation enthalpies for the successive elements, in general, is much higher.

• Mostly IE1<IE2 <IE3 in each group

• The increase in IE is primarily due to increase in nuclear charge. As the transition elements involve the gradual filling of (n-1)d orbitals, the effect of increase in nuclear charge is partly cancelled by the increase in screening effect.

• Consequently, the increase in I.E along the periods of d – block elements is very small.

3d < 4d < 5d (in 5d series - ineffective shielding by 4f electrons)

Relation between I.E and Stability of a metal in a given oxdn state

• With the help of I.E, we can predict which of the two metals in a given oxdn state is thermodynamically more stable.

Eg• When a metal M (0) is converted into M(11),

the energy required is equal to I1 + I2

Similerly M (IV) = I1 + I2+ I3 + I4

• Ni (0) Ni (II) I1 + I2 =2.49 x 103 kJ mol -1

• Pt (0) Pt (II) I1 + I2 =2.66 x 103 kJ mol -1

• Ni (0) Ni (IV) I1 + I2+ I3 + I4 =11.299 x 103 kJ mol -1

• Pt (0) Pt (IV) I1 + I2+ I3 + I4 =9.36 x 103 kJ mol -1

I1 + I2 for Ni (II) is less than I1 + I2 for Pt (II). So Ni (II) is more stable

Similarly Pt (IV) is more stable

OXIDATION STATES

+3

• One of the notable features of a transition element is the great variety of oxidation states it may show in its compounds

• Stability of a particular oxdn state depends up on nature of the element with which the transition metals form the compound

• The elements which give the greatest number of oxidation states occur in or near the middle of the series. Manganese, for example, exhibits all the oxidation states from +2 to +7.

• Elements in the beginning of the series exhibit fewer oxidation state (have small no. of electrons in which they lose or contribute for sharing).

• Elements at the end of the series shows fewer oxdn states because they have too many electrons in d – orbitals. So they have few vacant d – orbitals which can be invoved in bonding.

• Lower oxdn state – Covalent character• Higher oxdn state – ionic• Higher oxdn states are more stable for heavier

members.Eg : in group VI, Mo (VI) and W (VI) are more stable

than Cr (VI). So Cr (VI) act as strong oxidizing agent.• The highest oxdn state - +8 (Ruthenium and

Osmium). • Low oxidation states are found when a complex

compound has ligands capable of π-acceptor character in addition to the σ-bonding. For example, in Ni(CO)4 and Fe(CO)5, the oxidation state of nickel and iron is zero.

Trends in Stability of Higher Oxidation States

• Stability – compounds with F and Oxygen• The ability of Fluorine to stabilize the highest

oxidation state is due to either high lattice energy as in case of CoF3 or high bond enthalpy as in case of VF5 and CrF6.

• The ability of Oxygen to stabilize the highest oxidation state is due to its ability to form multiple bonds with metals.

Stable halides of first transition elementsOxdn no.

4 5 6 7 8 9 10 11 12

+6 Cr F6

+5 VF5 Cr F5

+4 TiX4 VX4 I Cr X4 MNf4

+3 TiX3 VX3 Cr X3 MnF3 Fe X3 Co F3

+2 TiX2 III VX2I Cr X2 MnX2 Fe X2 Co X2 Ni X2 Cu X2

II ZnX2

+1 Cu XIII

X = F to I, XII = F, XI = F to Br , X III = Cl to I

• The highest oxidation numbers are achieved in TiX4 (tetrahalides), VF5 and CrF6. The +7 state for Mn is not represented in simple halides but MnO3F is known, and beyond Mn, no metal has a trihalide except FeX3 and CoF3.

• Although V(V )is represented only by VF5, the other halides, however, undergo hydrolysis to give oxohalides, VOX3. Another feature of fluorides is their instability in the low oxidation states e.g., VX2 (X = CI, Br or I)

• All Cu(II) halides are known except the iodide. In this case, Cu2+ oxidises I– to I2:

2Cu2+ + 4I- → Cu2I2 (s) + I 2

• However, many copper (I) compounds are unstable in aqueous solution and undergo disproportionation.

2Cu2+ → Cu2+ + Cu• The stability of Cu2+ (aq) rather than Cu+(aq) is

due to the much more negative ΔhydH0 of Cu2+

(aq) than Cu+, which more than compensates for the second ionisation enthalpy of Cu.

• Transition metals also exhibits the highest Oxdn state in their Oxides.

• The ability of Oxygen to stabilize higher oxidation states are much higher than Fluorine..

• The highest Oxdn state with Fluorine by Mn is +4 in MnF4 while it is + 7 in Mn2O7.

• Oxygen has the ability to form Multiple bonds with Metal atom.

The oxides of 3 – d transition elements are given below :

Oxdn No

3 4 5 6 7 8 9 10 11 12

+7 Mn2O7

+6 CrO3

+5 V2O5 MnO2

+4 TiO2 V2O4 CrO2 Mn2O3 Fe2O3

+3 Sc2O3 Ti2O3 V2O3Cr2O3 Mn3O4 Fe3O4 Co3O4

+2 TiO VO CrO MnO FeO CoO NiO CuO ZnO

+1 Cu2O

• The highest oxidation number in the oxides coincides with the group number and is attained in Sc2O3 to Mn2O7.

• Beyond Group 7, no higher oxides of Fe above Fe2O3, are known, although ferrates (VI) (FeO4)2–, are formed in alkaline media but they readily decompose to Fe2O3 and O2.

• Besides the oxides, oxocations stabilise V(v) as VO2

+, V(IV) as VO2+ and Ti(IV) as TiO2+.

STANDARD ELECTRODE POTENTIAL

• ELECTRODE POTENTIALS ARE THE MEASURE OF THE VALUE OF TOTAL ENTHALPY CHANGE.

• Electrode Potentials value depends enthalpy of atomization ΔHa & hydration ΔH hyd

• Lower the std E. P (Eo red), the more stable is the oxdn state of the metal in aqueous state.

The E0(M2+/M) value for copper is positive (+0.34V) : high ΔHa and low ΔH hyd). --- GREATER AMNT OF ENERGY REQUIRED TO TRANSFORM Cu INTO Cu2+

• Due to +ve Eo, Cu does not liberate hydrogen from acids.

• The general trend towards less negative Eo

values across the series is related to the general increase in the sum of the first and second ionisation enthalpies.

• It is interesting to note that the value of Eo for Mn, Ni and Zn are more negative than expected from the trend.

• The stability of the half-filled d sub-shell in Mn2+

and the completely filled d10 configuration in Zn2+

are related to their Eo values, whereas Eo for Ni is related to the highest negative ΔhydHo.

• The low value for Sc reflects the stability of Sc3+

which has a noble gas configuration. The highest value for Zn is due to the removal of an electron from the stable d10 configuration of Zn2+. The comparatively high value for Mn shows that Mn2+

(d5) is particularly stable, whereas comparatively low value for Fe shows the extra stability of Fe3+

(d5).

CHEMICAL REACTIVITY• Transition metals vary widely in their chemical

reactivity. Many of them are sufficiently electropositive to dissolve in mineral acids, although a few are ‘noble’—that is, they are unaffected by simple acids.

• The metals of the first series with the exception of copper are relatively more reactive and are oxidised by 1M H+, though the actual rate at which these metals react with oxidising agents like hydrogen ion (H+) is sometimes slow.

• The EO valuesfor M2+/M indicate a decreasing tendency to form divalentcations across the series.

• This general trend towards less negative EO values is related to the increase in the sum of the first and second ionisation enthalpies.

• It is interesting to note that the EO values for Mn, Ni and Zn are more negative than expected from the general trend.

• EO values for the redox couple M3+/M2+ shows that Mn3+ and Co3+ ions are the strongest oxidising agents in aqueous solutions. The ions Ti2+, V2+ and Cr2+ are strong reducing agents and will liberate hydrogen from a dilute acid,

e.g.,• 2 Cr2+(aq) + 2 H+(aq) → 2 Cr3+(aq) + H2(g)

MAGNETIC PROPERTIES• Substances which contain species

(Atoms/ions/molecules) with unpared electrons in their orbitals – PARAMAGNETIC.

• PARAMAGNETIC SUBSTANCES are weakly attracted by the magnetic field.

• Strongly attracted called FERROMAGNETIC.• Substances which do not contain any unpaired

electrons and are repelled my magnetic field _ DIAMAGNETIC.

• Transition metals usually contains unpaired electrons – so it is paramagnetic.

• Paramagnetic behavior increases with increase in unpaired electron.

• Paramagnetism expressed in terms of Magnetic moment., it is related to no. of unpaired electrons.

• The magnetic moments calculated from the ‘spin-only’ formula and those derived experimentally.

Magnetic moment µ = √ n(n+2) BM

n- no. of unpaired electronsBM – Bohr magnetone (unit of M.M) BM = 9.27x10-21 erg/gauss• Single unpaired electronhas a magnetic

moment of 1.73 Bohr magnetons (BM).• magnetic moment of an electron is due to its

spin angular momentum and orbital angular momentum

Formation of Coloured Ions• When an electron from a lower energy d

orbital is excited to a higher energy d orbital, the energy of excitation corresponds to the frequency of light absorbed.

• This frequency generally lies in the visible region. The colour observed corresponds to the complementary colour of the light absorbed.

• The frequency of the light absorbed is determined by the nature of the ligand.

• Zn 2+ / Cd 2+ - all d orbitals are fully filled• Ti 4+ - all d orbitals are vacant so, no d – d transition occurs. Therefor they

do not absorb radiations. So they are colourless.

Formationof Complex Compounds•Metal ions bind a number of anions or neutral molecules giving complex[Fe(CN)6]3–, [Fe(CN)6]4–, [Cu(NH3)4]2+ and [PtCl4]2–

.This is due to the •Comparatively smaller sizes of the metal ions,

• Their high ionic charges and

•The availability of d orbitals for bond formation.

Formation of Interstitial Compounds

•When small atoms like H, C or N are trapped inside the crystal lattices of metals

•They are usually non stoichiometric

•example, TiC, Mn4N, Fe3H, VH0.56 and TiH1.7

(i) They have high melting points, higher than those of pure metals.(ii) They are very hard, some borides approach diamond in hardness.(iii) They retain metallic conductivity.(iv) They are chemically inert.

Alloy Formation

• Because of similar radii and other characteristics of transition metals,

• The alloys so formed are hard and have often high melting points.

• ferrous alloys: chromium, vanadium, tungsten, molybdenum and manganese are used for the production of a variety of steels and stainless steel.

• Alloys of transition metals with non transition metals such as brass (copper-zinc) and bronze (copper-tin),

CATALYTIC ACTIVITY• The transition metals and their compounds

are known for their catalytic activity. • This activity is ascribed to their ability to adopt

multiple oxidation states and to form complexes.

DISPROPORTIONATION• When a particular oxidation state becomes less

stable relative to other oxidation states, one lower, one higher, it is said to undergo disproportionation. For example, manganese (VI) becomes unstable relative to manganese(VII) and manganese (IV) in acidic solution.

3 MnVIO4 2– + 4 H+ → 2 MnVIIO–4 + MnIVO2 + 2H2O

Oxides and Oxoanions of Metals

• The elements of first transition series form variety of oxides of different oxidation states having general formula MO, M2O3, M3O6, MO2, MO3.

• Theses oxides are generally formed by heating the metal with oxygen at high temperature.

Sc – Sc2O3 Basic

Ti – TiO Basic, Ti2O2 Basic, TiO2 Amphoteric

V – VO Basic, V2O3 Basic, VO2 Ampho, V2O5 Acidic

Cr – CrO Basic, Cr2O3 Ampho, CrO2 Ampho,

CrO3Acidic

Mn – MnO basic, Mn2O3 Basic, Mn3O4 Ampho,

MnO2 Ampho, Mn2O7 Acidic

Fe – FeO Basic, Fe2O3 Amph, Fe3O4 Basic

Co – CoO Basic Ni – NiO BasicCu – Cu2O Basic, CuO Ampho

Zn – ZnO Ampho

• In general lower oxidation state metal – BASIC Higher oxidation state metal – ACIDIC Intermediate oxidation state - AMPHOTERIC• ExampleMnO (+2)basic, Mn2O3 (+3)Basic, Mn3O4 (+

8/3)Ampho, MnO2 (+4) Ampho, Mn2O7 (+7)Acidic

• The highest oxidation number in the oxides coincides with the group number and is attained in Sc2O3 to Mn2O7.

• Beyond Group 7, no higher oxides of Fe above Fe2O3, are known, although ferrates (VI) (FeO4)2–, are formed in alkaline media but they readily decompose to Fe2O3 and O2.

• Besides the oxides, oxocations stabilise V(v) as VO2

+, V(IV) as VO2+ and Ti(IV) as TiO2+.

• As the oxidation number of a metal increases, ionic character decreases. In the case of Mn, Mn2O7 is a covalent green oil. Even CrO3 and V2O5 have low melting points. In these higher oxides, the acidic character is predominant.

Potassium dichromate K2Cr2O7

STEP 1• Dichromates are generally prepared from

chromate which in turn are obtained by the fusion of chromite ore (FeCr2O4) with sodium or potassium carbonate in free access of air. The reaction with sodium carbonate occurs as follows:4 FeCr2O4 + 8 Na2CO3 + 7 O2 → 8 Na2CrO4+2

Fe2O3 + 8 CO2

STEP 2• The yellow solution of sodium chromate is

filtered and acidified with sulphuric acid to give a solution from which orange sodium dichromate, Na2Cr2O7. 2H2O can be crystallised.

2Na2CrO4 + H2SO4 → Na2Cr2O7 + Na2SO4 + H2O

STEP 3Conversion of Sodium dichromate in to Potassium

dichromateNa2Cr2O7 + 2 KCl → K2Cr2O7 + 2 NaCl

• The oxidation state of chromiumin chromate and dichromate is the same.

2 CrO42– + 2H+ → Cr2O7

2– + H2O

Cr2O72– + 2 OH- → 2 CrO4

2– + H2O• The chromate ion is tetrahedral whereas the

dichromate ion consists of two tetrahedra sharing one corner with Cr–O–Cr bond angle of 126°.

•Sodium and potassium dichromates are strong oxidising agentsPotassium dichromate is used as a primary standard in volumetric analysis. In acidic solution, its oxidising action can be represented as follows:

Cr2O72– + 14H+ + 6e– → 2Cr3+ + 7H2O (EV = 1.33V)

• acidified potassium dichromate will oxidise iodides to iodine, sulphides to sulphur, tin(II) to tin(IV) and iron(II) salts to iron(III). The half-reactions are noted below:

• 6 I– → 3I2 + 6 e– ;

• 3 H2S → 6H+ + 3S + 6e–

• 3 Sn2+ → 3Sn4+ + 6 e–

• 6 Fe2+ → 6Fe3+ + 6 e–

Cr2O72– + 14 H+ + 6 Fe2+ → 2 Cr3+ + 6 Fe3+ + 7 H2O

Potassium permanganate KMnO4

• Potassium permanganate is prepared by fusion of MnO2 with an alkali metal hydroxide and an oxidising agent like KNO3. This produces the dark green K2MnO4 which disproportionates in a neutral or acidic solution to give permanganate.

2MnO2 + 4KOH + O2 → 2K2MnO4 + 2H2O

3MnO42– + 4H+ → 2MnO4

– + MnO2 + 2H2O

The manganate and permanganate ions aretetrahedral; the green manganate is paramagnetic with one unpaired electron but the permanganate is diamagnetic.

THE INNER TRANSITION ELEMENTS ( f-BLOCK)

• The elements in which the additional electrons enters (n-2)f orbitals are called inner transition elements. The valence shell electronic configuration of these elements can be represented as (n – 2)f0-14(n – 1)d0-1ns2.

• 4f inner transition metals are known as lanthanides because they come immediately after lanthanum and 5f inner transition metals are known as actinoids because they come immediately after actinium.

Electronic ConfigurationElement name Symbol Z Ln Ln3+ Radius

Ln3+/ pmLanthanum La 57 [Xe]6s25d1 [Xe]4f0 116Cerium Ce 58 [Xe]4f16s25d1 [Xe]4f1 114Praesodymium Pr 59 [Xe]4f36s2 [Xe]4f2 113Neodymium Nd 60 [Xe]4f46s2 [Xe]4f3 111Promethium Pm 61 [Xe]4f56s2 [Xe]4f4 109Samarium Sm 62 [Xe]4f66s2 [Xe]4f5 108Europium Eu 63 [Xe]4f76s2 [Xe]4f6 107Gadolinium Eu 64 [Xe]4f76s25d1 [Xe]4f7 105Terbium Tb 65 [Xe] 4f96s2 [Xe]4f8 104Dysprosium Dy 66 [Xe] 4f106s2 [Xe]4f9 103Holmium Ho 67 [Xe] 4f116s2 [Xe]4f10 102Erbium Er 68 [Xe] 4f126s2 [Xe]4f11 100Thulium Tm 69 [Xe] 4f136s2 [Xe]4f12 99Ytterbium Yb 70 [Xe] 4f146s2 [Xe]4f13 99Lutetium Lu 71 [Xe] 4f146s25d1 [Xe]4f14 98

Atomic and ionic sizes: The Lanthanide Contraction

• As the atomic number increases, each succeeding element contains one more electron in the 4f orbital and one proton in the nucleus. The 4f electrons are ineffective in screening the outer electrons from the nucleus causing imperfect shielding. As a result, there is a gradual increase in the nucleus attraction for the outer electrons. Consequently gradual decrease in size occur. This is called lanthanide contraction.

Consequences of L. C• There is close resemblance between 4d and

5d transition series.• Ionization energy of 5d transition series is

higher than 3d and 4d transition series.• Difficulty in separation of lanthanides

Ionization Enthalpies

• Fairly low I. E• First ionization enthalpy is around 600 kJ mol-1,

the second about 1200 kJ mol-1 comparable with those of calcium.

• Due to low I. E, lanthanides have high electropositive character

Coloured ions• Many of the lanthanoid ions are coloured in

both solid and in solution due to f – f transition since they have partially filled f – orbitals.

• Absorption bands are narrow, probably because of the excitation within f level.

• La3+ and Lu3+ ions do not show any colour due to vacant and fully filled f- orbitals.

Magnetic properties

• The lanthanoid ions other then the f 0 type (La3+ and Ce3+) and the f14 type (Yb2+ and Lu3+) are all paramagnetic. The paramagnetism rises to the maximum in neodymium.

• Lanthanides have very high magnetic susceptibilities due to their large numbers of unpaired f-electrons.

Oxidation States• Predominantly +3 oxidation state.• +3 oxidation state in La, Gd, Lu are especially

stable ( Empty half filled and Completely filled f – subshell respectively)

• Ce and Tb shows +4 oxdn state ( Ce 4+ - 4fo & Tb 4+ 4f7)

• Occasionally +2 and +4 ions in solution or in solid compounds are also obtained.

• This irregularity arises mainly from the extra stability of empty, half filled or filled f subshell.

• The most stable oxidation state of lanthanides is +3. Hence the ions in +2 oxidation state tend to change +3 state by loss of electron acting as reducing agents whereas those in +4 oxidation state tend to change to +3 oxidation state by gain of electron acting as a good oxidising agent in aqueous solution.

• Why Sm2+, Eu2+, and Yb2+ ions in solutions are good reducing agents but an aqueous solution of Ce4+ is a good oxidizing agent?

properties

• Silvery white soft metals, tarnish in air rapidly• Hardness increases with increasing atomic

number, samarium being steel hard.• Good conductor of heat and electricity.• Promethium - Radioactive

Chemical Properties • Metal combines with hydrogen when gently

heated in the gas. • The carbides, Ln3C, Ln2C3 and LnC2 are formed

when the metals are heated with carbon.• They liberate hydrogen from dilute acids and

burn in halogens to form halides.• They form oxides and hydroxides, M2O3 and

M(OH)3, basic like alkaline earth metal oxides and hydroxides.

Ln W ith helogensHeated w ith S

C 2773 K

Ln S2 3

2

32 2LnN LnC Ln(OH) +H

3LnX

HLn O2 3

The Actinides

• All isotopes are radioactive, with only 232Th, 235U, 238U and 244Pu having long half-lives.

• Only Th and U occur naturally-both are more abundant in the earth’s crust than tin.

• The others must be made by nuclear processes.

• The dominant oxidation state of actinides is +3. Actinides also exhibit an oxidation state of +4. Some actinides such as uranium, neptunium and plutonium also exhibit an oxidation state of +6.

• The actinides show actinide contraction (like lanthanide contraction) due to poor shielding of the nuclear charge by 5f electrons.

• All the actinides are radioactive. Actinides are radioactive in nature.

Actinoide Contraction

• The size of atoms / M3+ ions decreases regularly along actinoid seris. The steady decrease in ionic/ atomic radii with increase in atomic number is called Actinoide Contraction.

• The contraction is greater from element to element in this series – due to poor shielding effect by 5 f electron.

Electronic configurationElement name Symbol Z Ln Ln3+ Radius

Ln3+/ pmActinium Ac 89 [Rn] 6d17s2 [Rn]4f0 111Thorium Th 90 [Rn ]5d27s2 [Rn]4f1 Protactinium Pa 91 [Rn]5f26d17s2 [Rn]4f2 Uranium U 92 [Rn]5f36d17s2 [Rn]4f3 103Neptunium Np 93 [Rn]5f46d17s2 [Rn]4f4 101Plutonium Pu 94 [Rn]5f67s2 [Rn]4f5 100Americium Am 95 [Rn]5f77s2 [Rn]4f6 99Curium Cm 96 [Rn]5f76d17s2 [Rn]4f7 99Berkelium Bk 97 [Rn]5f97s2 [Rn]4f8 98Californium Cf 98 [Rn]5f107s2 [Rn]4f9 98Einsteinium Es 99 [Rn]5f117s2 [Rn]4f10 Fermium Fm 100 [Rn]5f127s2 [Rn]4f11 Mendelevium Md 101 [Rn]5f137s2 [Rn]4f12 Nobelium No 102 [Rn]5f147s2 [Rn]4f13 Lawrencium Lr 103 [Rn]5f146d17s2 [Rn]4f14

Magnetic properties• Paramagnetic behaviour• Magnetic properties are more complex than

those of lanthanoids.

M.P and B.PHigh M.P and B.PDo not follow regular gradation of M.P or B.P

with increase in atomic number

IONISATION ENTHALPY• Low I.E. so electropositiity is HighCOLOUR• Generally coloured• Colour depends up on the number of 5 f

electrons• The ions containing 5 f o and 5 f 7 are

colouressEg – U 3+ (5 f 3 ) – RedNP 3+ (5 f 4 ) – Bluish

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